CHARGED PARTICLES IN MATTER The presence of charged particles in matter is suggested by the phenomena of static electricity and electricity conduction through certain substances. Therefore, Atoms can be divided further into sub-atomic particles. *SUB-ATOMIC PARTICLES An atom is the smallest unit of an element that still retains its chemical properties. It consists of three fundamental sub-atomic particles: Electrons (e−): These are the negatively charged particles having very little mass. They orbit the nucleus. Protons (p+): These are the positively charged particles. They are located in the nucleus and have a mass of about one atomic mass unit. Neutrons (n): These are the particles that are neutral (have no charge). They are also found in the nucleus and have a mass roughly equivalent to that of a proton. EARLY MODELS OF ATOM *DALTON’S ATOMIC THEORY Dalton’s theory states that matter is made up of small indestructible atoms that combine in set ratios and rearrange in reactions without being created or destroyed. This theory suggested that Atom is indivisible – which could not be broken down into smaller particles. # But the discovery of Sub-Atomic Particles inside the atom disproved this principle of Dalton’s atomic theory. *DISCOVERY OF ELECTRONS (1897) -Given by: J.J. Thomson in 1897. -J.J. Thomson discovered the electron in 1897 through the cathode ray experiment, proving atoms are not indivisible. -Experiment Setup -Key Observations -Conclusion -Thomson’s Contribution -Characteristics of an Electron *DISCOVERY OF PROTON -Given by: E. Goldstein in 1886 –E. Goldstein discovered protons in 1886 through canal rays, identifying them as positively charged particles inside atoms. -Experiment Setup -Key Observations -Conclusion -Importance -characteristics of proton *DISCOVERY OF NEUTRON – Given by: James Chadwick -In 1932, James Chadwick discovered the neutron, a neutral particle in the nucleus with mass equal to a proton. -Experiment Key Observations -Conclusion -Importance -Characteristics of Neutron ATOMIC MODEL There are three Atomic Models on Arrangement of Sub-Atomic Particles. 1) Thomson’s Model of the Atom (1898) -After discovering the electron, J.J. Thomson proposed a model to explain the structure of the atom. -Thomson imagined the atom as a sphere of positive charge with electrons studded inside it, called the Plum Pudding Model. -Main Features -Limitations -Importance 2) Rutherford’s Model (Nuclear Model) -Rutherford performed the Gold Foil Experiment (α-particle scattering experiment) with Geiger and Marsden. -Aim: To test Thomson’s Plum Pudding Model. -Experiment Setup -Observations -Main Features of Rutherford’s Model -Limitations -Importance 3) Bohr’s Model -Niels Bohr improved Rutherford’s model. -Solved the stability problem and explained hydrogen spectrum. -Postulates of Bohr’s Model -Achievements –Limitations Atomic Number and Mass Number An atom is made up of subatomic particles: protons, neutrons, and electrons. The atomic number and mass number are two important terms used to describe the composition of an atom. *Atomic Number (Z) The atomic number, symbolized by Z, is the number of protons in the nucleus of an atom. Formula: Atomic Number(Z)=Number of Protons For a neutral atom: Number of Electrons=Number of Protons *Mass Number (A) The mass number, symbolized by A, is the total number of protons and neutrons in the nucleus of an atom. Since the mass of electrons is negligible, the mass number essentially represents the total mass of the atom’s nucleus. Formula: Mass Number(A)=Number of Protons + Number of Neutrons You can rearrange this formula to find the number of neutrons: Number of Neutrons = Mass Number(A) – Atomic Number(Z) *Isotope Notation Elements are often represented with their atomic and mass numbers. The notation is as follows: Where: Example: Carbon (C) A common isotope of carbon has an atomic number of 6 and a mass number of 12. This can be written as: Electron Distribution in Orbits After the discovery of protons, electrons, and neutrons, the next task was to understand how electrons are arranged within an atom. -Niels Bohr and Bury developed laws for the distribution of electrons in different shells (orbits/energy levels) surrounding the nucleus. *Rules for Electron Distribution 1. Naming of Shells 2. Maximum Number of Electrons in a Shell 3. Octet Rule (Stability Rule) 4. Filling of Electrons in Successive Shells *Importance of Electron Distribution VALENCY Valency is the combining capacity of an atom. It is the number of electrons an atom must gain, lose, or share to achieve a stable electronic configuration, typically having a full outermost shell (an octet of 8 electrons, or a duplet of 2 for elements like Helium). This stable state resembles the electronic configuration of a noble gas. -Atoms are stable when they have 8 electrons in their outermost shell (Octet Rule). -To become stable, atoms: *How to Determine Valency The valency of an element is determined by the number of electrons in its outermost shell, known as valence electrons. *Types of Valency Based on how atoms achieve stability, we can distinguish between two types of valency: *Noble Gases: Zero Valency Noble gases (Group 18 elements) like Helium, Neon, and Argon have a completely filled outermost shell (a stable octet or duplet). Because they are already stable, they do not need to gain, lose, or share electrons. #Valency of noble gases is zero. They are generally unreactive. *Importance of Valency ISOTOPES AND ISOBARS *Isotopes -Definition – Characteristics – Example: Chlorine has two isotopes, Chlorine-35 and Chlorine-37. – Uses of Isotopes *Isobars – Definition – Characteristics
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